orbitaloverlapping

Orbital overlap refers to the region in space where atomic orbitals of two atoms intersect and interact to form a chemical bond. In valence bond theory, bonding arises when orbitals overlap to place electron density between nuclei, which lowers the system's energy. The strength and length of a bond correlate with the extent of overlap between the participating orbitals.

Two primary modes are end-on (sigma) overlap, along the internuclear axis, and side-by-side (pi) overlap of adjacent p orbitals. Sigma bonds, from s-s, s-p, or p-p overlaps, are typically stronger because electron density lies directly between nuclei. Pi bonds arise from lateral overlap of parallel p orbitals above and below the bond axis and are usually weaker; they occur in multiple bonds. In transition metals, d orbital overlaps can contribute to delta interactions, though this is more specialized.

Overlap is often described quantitatively by the overlap integral S, which measures the extent to which two orbital wavefunctions coincide. A larger S leads to greater electron density between nuclei and a stronger, shorter bond. The energy match and phase alignment of the orbitals also determine whether the interaction is bonding or antibonding.

Factors that influence overlap include the sizes and energies of the interacting orbitals, their orientation and symmetry, and the internuclear distance. The formation of hybrid orbitals in valence bond theory (such as sp, sp2, sp3) is often described as mixing orbitals to maximize overlap with available orbitals on the neighboring atom.

In molecular orbital theory, overlap is treated more globally: atomic orbitals combine to form molecular orbitals that extend over the entire molecule, with bonding and antibonding combinations determined by symmetry and energy. Together, the concept of orbital overlap explains the directional character and strength of chemical bonds and is a foundational idea in understanding bonding and reactivity.